In 1785, researchers working with early, arcing electrical discharge devices had noted ozone’s peculiar odor (Christie 195). The word ozone is in fact derived from the Greek ozein, meaning “to smell.” In 1872, scientists determined that ozone was a triatomic form of oxygen, or O3. As a concentrated gas, ozone is pale blue, carries a strong odor, and is highly poisonous.
Spacially, ozone’s three oxygen atoms form an enormous molecular triangle. Because of weak atomic bonding between the distant atoms, ozone is a very unstable molecule that quickly dissociates into common oxygen and a free oxygen atom (O) called atomic oxygen.
The ability to release atomic oxygen makes ozone a powerful oxidizing agent, or giver of oxygen to other molecules. In contrast, common oxygen is two oxygen atoms held together by strong bonds that produce the stable O2 molecule. Common oxygen makes up 20 percent of our atmosphere.
Ozone forms naturally in the lower levels of the atmosphere when the electrical arcs of lightning pass through oxygen molecules.
After it is formed in the atmosphere, ozone survives only about 20 minutes before it dissociates. Most of Earth’s natural ozone, however, is formed in the stratosphere, the upper portion of the atmosphere. There, at an altitude of about 15 miles, photochemical reactions driven by intense solar ultraviolet radiation both create and destroy ozone.
High-energy ultraviolet rays split common oxygen molecules into free oxygen atoms, some of which combine with common oxygen molecules to form ozone. Simultaneously, low-energy ultraviolet radiation splits ozone into common oxygen and free oxygen atoms. Each free oxygen atom may then combine with common oxygen to create more ozone, or it may join with another oxygen atom to form common oxygen. Ozone can also react with nitrogen and hydrogen, and as well with chlorine, trace amounts of which originate naturally in soils, oceans, and volcanic eruptions and migrate into the stratosphere. Given the multitude of factors, especially the solar cycle and variations in stratospheric winds affecting the formation and lifetime of stratospheric ozone, its concentration and extent varies periodically.
The hole in the ozone layer
In 1970, a British research group working in Antarctica discovered an unusual reduction in stratospheric ozone (Christie 38). Over the next decade, studies revealed that ozone was being destroyed by a number of chemicals, primarily trace amounts of man-made chlorofluorocarbons (CFSs) – complex compounds of chlorine, fluorine, and carbon widely used in industry and manufacturing (Christie 29). CFCs are stable molecules that are unreactive in the lower atmosphere. But when these same molecules migrate in small quantities into the stratosphere, the intense ultraviolet radiation there splits them, releasing chlorine atoms that quickly react with ozone to form chlorine monoxide (ClO) and common oxygen. In turn, chlorine monoxide contacts free oxygen atoms, reacting to form common oxygen and a free chlorine atom – which then destroys another ozone molecule. Scientists estimate that in repetitive reactions a single chlorine atom can destroy more than 10,000 ozone molecules before it finally locks up in a less reactive form.
Since the 1950s, nontoxic, nonflammable, and inexpensive CFCs have made ideal aerosol propellants, refrigerants, and solvents. Aerosol cans, which are designed to release their entire contents into the air, accounted for most of the CFCs emitted into the atmosphere. In 1973 alone, aerosol manufacturers filled nearly six billion aerosol cans worldwide. Half of those employed compressed CFCs to propel the active ingredients – everything from hair spray and underarm deodorant to insecticide, paint, polish, and disinfectant – out of the can. By 1975, growing evidence of ozone layer depletion triggered a public-relations battle between aerosol can manufacturers and environmentalists that became known as “the Spray-Can War” (Christie 21).
Concern about the ozone layer spurred unprecedented international environmental cooperation. In 1987 in Montreal, 36 nations met to ratify the Vienna Agreement for the Protection of the Ozone Layer. The so-called Montreal Protocol banned production of the most widely used CFCs by 1995 and currently mandates termination of all production by 2010 (Fishman 45). Today more than 100 nations have additionally agreed to phase out production of methyl bromide, a pesticide that also harms the ozone layer.
By the late 1970s it could be shown that CO reacts with OH in the atmosphere much faster than had previously been believed (Christie 79). Furthermore, it is likely that a by-product of this atmospheric photochemical process is the production of ozone in the troposphere. The first reaction proceeds:
CO + OH → CO 2 + H
The H atom then very quickly (within a small fraction of a second) latches onto an oxygen molecule to form a reactive peroxy radical, HO 2. This HO 2 radical is analogous to the RO 2 radicals produced in the urban smog scenario discussed earlier. The key to whether or not ozone is produced in the remote atmosphere centers on what happens to the HO 2 radical. If there is a molecule of nitric oxide (NO) around, then it is likely that the following reaction takes place in the atmosphere:
HO 2 + NO → NO 2 + OH
The nitrogen dioxide will be photolyzed by visible light and ozone will again be made, just as it is in the polluted urban environment:
NO 2 + (visible) photon → NO + O
O + O 2 + M → O 3 + M
The above sequence of reactions that results in ozone formation is a catalytic cycle with respect to the hydroxyl radical and nitric oxide. Both OH and NO are returned to the atmosphere so that more carbon monoxide can be oxidized to carbon dioxide by the hydroxyl radical, and likewise so that nitric oxide can be converted to nitrogen dioxide to make more ozone (Fishman 90).
Any molecule can be broken into smaller fragments if it absorbs energy from sufficiently short wavelength ultraviolet radiation. In the case of CFCs, the ozone layer filters out all of the wavelengths that might break up the molecules. But if they were to travel to 15 km altitude and higher, they would start to rise above some of the ozone. Then some of the ultraviolet light that could break the molecules down into smaller fragments would not be so effectively blocked. The molecules would be broken into very reactive free radicals by any of this light that got through. ultraviolet light with wavelength between 190 and 215 nm can decompose CFC molecules, splitting off atomic chlorine (Christie 78).
Any CFC molecules that found their way into the stratosphere would encounter this light, and be transformed from unreactive materials to very reactive chlorine atoms. Studies of ultraviolet (UV) radiation and its absorption also contributed to early understanding of ozone. The spectrum of sunlight reaching the earth showed a sharp cutoff in the UV region of wavelengths shorter than about 290 nanometers (nm), suggesting that some atmospheric substance might be absorbing radiation with shorter wavelengths (Fishman 7). Hartley measured ozone’ absorption and found that even tiny quantities absorbed strongly in the 200–300 nm region of the UV, with the edge of the strong absorption region corresponding closely to the sharp cutoff in sunlight. Moreover, using an artificial source of UV light showed that surface air readily transmitted wavelengths down to 250 nm, well below the cutoff in incoming sunlight.
Oxygen molecule and atoms
Oxygen molecules have two atoms (O 2), the ozone molecule has three oxygen atoms (O 3). The process that makes ozone on the surface is similar to the one producing it in the stratosphere, with one important difference. At the surface there aren’t enough high-energy photons to break down the oxygen molecule. Still, there are other gases that can be used to create ozone. One of them is nitrogen dioxide (NO 2) (Christie 83). NO 2 has an affinity for the less energetic photons in the visible portion of the electromagnetic spectrum (i.e., sunlight). These low-energy photons break down NO 2 into NO (nitric oxide), plus a free oxygen atom. It is these free oxygen atoms, released from bondage to the nitrogen dioxide molecule, that then team up with oxygen to form ozone.
As this photon approaches the earth, the chances for it to hit a molecule in the earth’s atmosphere become considerably greater as the air density increases closer to the surface. If the photon hits an oxygen molecule, which is comprised of two oxygen atoms, it will photolyze this molecule into its two atoms. The process can be described by the chemical equation:
O 2 + (high energy) photon → O + O
Once the oxygen atoms are “free,” they can recombine with an oxygen molecule to form ozone (O 3). But if they are to do this, a third neutral molecule must show up to absorb the excess energy the oxygen atoms have recently acquired after being “zapped” by the high-energy photon. In most cases this third molecule is nitrogen (N 2), but it can be either another oxygen molecule or an argon (Ar) atom. Chemically, the process is described:
O + O 2 (+M) → O 3 (+M)
where M is either N 2, O 2, or Ar.
So at the very high levels of the atmosphere, where there is relatively little N 2, O 2, or Ar present, there is much more atomic oxygen (O) than ozone (O 3) (Fishman 37). As we get lower in the atmosphere, O 3 becomes more abundant, relative to O. Again, however, we should note that atomic oxygen and ozone always remain trace constituents of the atmosphere, never reaching concentrations above the part per million range.
Importance to Life
Any significant depletion of the earth’s ozone shield over temperate latitudes would be a major disaster for humanity, not to mention the rest of the planet’s inhabitants. Recent studies suggest that for every 1 percent drop in the amount of ozone above a given locality, there will be a 2 to 3 percent increase in the annual number of malignant melanoma skin cancers. Last year about 23,000 Americans were diagnosed as having a malignant melanoma, and 6,000 died from it.
The ozone created up in the stratosphere is called the ozone shield, which is chiefly responsible for permitting life as we know it to exist here on the surface of the earth. Without the ozone interacting with the high-energy photons of the ultraviolet portion of the sun’s spectrum, those same photons would reach the earth, bombarding us with more of the harmful ultraviolet rays of the sun’s spectrum. If these harmful ultraviolet rays were allowed to penetrate to the surface, many of the life forms on the planet would not survive since the delicate balance that supports such life would be destroyed. Ozone’s importance to life stems from its ability to block the sun’s deadly ultraviolet-B light. Each 1 percent reduction in the amount of atmospheric ozone allows 2 percent more UV-B radiation to reach the earth’s surface (Christie 200).
Distribution of Ozone in the Stratosphere
With the measurement of the reaction rates in the laboratory, the picture of stratospheric chemistry became much more complicated than had been previously explained by Chapman’s pure oxygen chemistry (Parson 70). Furthermore, the inclusion of these radicals in a set of calculations that were being used to describe the distribution of ozone in the stratosphere yielded a computed distribution that was more consistent with the growing data base of ozone measurements. However, it is easy to see how the science of atmospheric chemistry became considerably more complicated once additional molecules had to be considered to explain the distribution of ozone in the stratosphere.
A further refinement of stratospheric chemistry came about in 1970 when Paul Crutzen, then of the University of Stockholm, and Harold Johnston, of the University of California at Berkeley, independently introduced theories proposing that molecules that contain nitrogen, analogous to those that contain hydrogen, play an important role in the chemistry of the stratosphere. Their proposals supported a growing concern at the time that direct injection of nitrogen oxides by high-flying aircraft might seriously impact the chemistry of the ozone layer.
The world organized to determine and quantify the extent to which the ozone in the stratosphere had been depleted. Furthermore, if a decrease of ozone in the stratosphere could be definitively established, this panel was charged with the responsibility of finding out whether the decrease could be attributed to natural or anthropogenic (human-produced) causes. Prior to the establishment of this fact-finding panel, several scientific papers had been published which reported that ozone amounts in the stratosphere had been decreasing in the late 1970s and early 1980s. There were, however, several possible explanations for the observed decreasing trend.
In 1973, scientists discovered that human-produced chlorofluorocarbons (CFCs) decimate ozone when they break apart six to 25 miles above the Earth’s surface (Parson 225). This ozone destruction allows the more biologically harmful ultraviolet-B rays (UV-B) to reach the Earth’s surface. Every one percent decrease in ozone causes a 2.2 percent increase in DNA-damaging UV-B radiation. UV-B is also known to be the primary cause of human skin cancer and is linked to cataracts and immune system suppression. It also damages crops and marine algae. The best-known depletion – the ozone hole above Antarctica -measured nine million square miles in September 1992, a 15 percent increase over 1991 (Parson 137). Less well known is the steady depletion of ozone over populated areas worldwide. Australia, where frog species are in serious decline, is one area where the ozone layer has been most severly diminished.